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Recap Resonance

Recap Resonance. Elements from the third period and beyond form compounds with 8 or more electrons around the atom. This flexibility may result in quite different resonance structure being possible.

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Recap Resonance

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  1. Recap Resonance • Elements from the third period and beyond form compounds with 8 or more electrons around the atom. • This flexibility may result in quite different resonance structure being possible. • The resonance structure(s) with the greatest contribution to the actual structure can be identified using the valency of oxygen as a guide. Example: ClO4-

  2. VSEPR Theory Lewis structures give bonding arrangements but do not imply any molecular shape. For this we use: Valence Shell Electron Pair Repulsion Theory This relies on minimising repulsion between areas of electrons (bond pairs and lone pairs) around the central atom.

  3. VSEPR Theory • Draw Lewis Structure. • Count number of electron pairs. • Count both bonding pairs and non-bonding pairs. • Count multiple bonds as only one area of electrons. • Determine the arrangement of electron pairs. • Electron pairs want to be as far away from each other as possible. • Use atom positions to name molecular geometry. • This is the atom positions.

  4. Electron Pair Arrangements • Two electron pairs: • Atoms at the opposite ends of a line. • 180 degrees between areas of electrons. • Called linear. • eg CO2

  5. Electron Pair Arrangements • Three electron pairs: • Atoms at the corners of a triangle. • 120 degrees between electron pairs. • Called trigonal planar. • Eg BF3

  6. Electron Pair Arrangements • Four electron pairs: • Atoms at the corners of a tetrahedron. • 109.5 degrees between electron pairs. • Called tetrahedral. • Eg CH4

  7. Molecular Geometry Remove one arm from the electron pair arrangement for each lone pair present. Trigonal Planar (3 e- pairs) Figure 10.4 Silberberg

  8. Molecular Geometry Tetrahedral (4 e- pairs) Figure 10.5 Silberberg

  9. 104.5 107 109.5 Molecular Geometry • Repulsion: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

  10. Molecular Geometry Figure 10.9 Silberberg

  11. Molecular Geometry - Example • Molecules with multiple bonds eg COCl2 total 24 e- ~120 3 areas of electrons about C, so trigonal planar arrangement of electrons No lone pairs so molecular geometry is also trigonal planar

  12. Molecular Geometry - Example • Cases when there is no single central atom • Just apply the VSEPR rules to each central atom in turn. ~120 3 areas of electrons about each C, so trigonal planar arrangement of electrons about each C

  13. Dipole Moments • Any bond between two different atoms will be polar. • A molecule has a permanent dipole moment if it contains polar bondsand it is not a symmetrical shape. • Note: Cations and anions are not polar – the overall charge overwhelms any local d+ vs d- effects.

  14. Dipole Moments • Polar molecules • Non-polar molecules H2O HF CHCl3 CCl4 CO2 N2

  15. Learning Outcomes: By the end of this lecture, you should: work out the number of bonding and non-bonding pairs from the Lewis structure of a molecule predict the distribution of these pairs around an atom predict and describe the molecular shape determine if a permanent dipole exists be able to complete the worksheet (if you haven’t already done so…)

  16. Questions to complete for next lecture: • Draw the shapes of the following molecules and ions and give approximate bond angles • BH3 • NH4+ • CS2 • CH2O • CH3Cl • H3O+

  17. Questions to complete for next lecture: • What are the approximate C-C-C bond angles in the two molecules below? • Are these molecules flat?

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